On the basis of the following $E^\circ$ values, the strongest oxidizing agent is :
$[Fe(CN)_6]^{4-} \rightarrow [Fe(CN)_6]^{3-} + e^-$; $E^\circ = -0.35$ V
$Fe^{2+} \rightarrow Fe^{3+} + e^-$; $E^\circ = -0.77$ V

The given values are oxidation potentials. Reverse the reactions to get the standard reduction potentials:
$[Fe(CN)_6]^{3-} + e^- \rightarrow [Fe(CN)_6]^{4-}$; $E^\circ = +0.35$ V
$Fe^{3+} + e^- \rightarrow Fe^{2+}$; $E^\circ = +0.77$ V
An oxidizing agent oxidises other species and gets itself reduced easily.
Higher the positive reduction potential, stronger is the oxidizing agent.
$Fe^{3+}$ has the higher reduction potential (+0.77 V) compared with $[Fe(CN)_6]^{3-}$ (+0.35 V).
$Fe^{2+}$ and $[Fe(CN)_6]^{4-}$ are the reduced forms; they act as reducing agents, not oxidizing agents.
So the strongest oxidizing agent is $Fe^{3+}$.